This formula explicitly spells out the relationship between the number of bonding electrons and their relationship to how many are formally “owned” by the atom. However, since the “number of bonding electrons divided by 2” term is also equal to the number of bonds surrounding the atom, here’s theshortcut formula:

Let’s do one last example. Let’s do CH3+ (with no lone pairs on carbon). It’s the orange one on the bottom row.

The number of valence electrons for carbon is 4

The number of non-bonded electrons is zero

The number of bonds around carbon is 3.

So formal charge = 4 – (0 +3) = 4 – 3 = +1

You can apply this formula to any atom you care to name.

Here is a chart for some simple molecules along the series B C N O . I hope beryllium and fluorine aren’t too offended that I skipped them, but they’re really not that interesting for the purposes of this table.

Note the interesting pattern in the geometries (highlighted in colour): BH4(–), CH4, and NH4(+) all have the same geometries, as do CH3(–), NH3, and OH3(+). Carbocation CH3(+) has the same electronic configuration (and geometry) as neutral borane, BH3. The familiar bent structure of water, H2O, is shared by the amide anion, NH2(–). These shared geometries are one of the interesting consequences of valence shell electron pair repulsion theory (VSEPR – pronounced “vesper“, just like “Favre” is pronounced “Farve”.)

The formal charge formula also works for double and triple bonds:

Here’s a question. Alkanes, alkenes, and alkynes are neutral, since there are four bonds and no unbonded electrons: 4 – [4+0] = 0. For what other values of [bonds + nonbonded electrons] will you also get a value of zero, and what might these structures look like? (You’ll meet some of these structures later in the course).

One final question – why do you think this is called “formal charge”?

Think about what the formal charge of BF4 would be. Negative charge on the boron. What’s the most electronegative element here? Fluoride, of course, with an electronegativity of 4.0, with boron clocking in at 2.0. Where do you think that negative charge really resides?

Well, it ain’t on boron. It’s actually spread out through the more electronegative fluoride ions, which become more electron-rich. So although the “formal” address of the negative charge is on boron, the electron density is actually spread out over the fluorides. In other words, in this case the formal charge bears no resemblance to reality.

Another reminder – 10 videos with solved examples of formal charge problems, right here (look at the very top of the page)

In CH3 i think FC on C should be -1 as carbon valency is 4 it has already bonded with 3 hydrogen atom one electron is left free on carbon to get bond with or share with one electron H hence, number of non bonded electrons lone pair of electrons is considered as 2. 4-(2+3) = -1.
In your case if we take 0 than valency of c is not satisfied.

I hope the post doesn’t get interpreted as “formal charges have no significance”. If it does I will have to change some of the wording.

What I mean to get across is that formal charges assigned to atoms do not *always* accurately depict electron density on that atom, and one has to be careful.

In other words, formal charge and electron density are two different things and they do not always overlap.

Formal charge is a book-keeping device, where we count electrons and assign a full charge to one or more of the atoms on a molecule or ion.
Electron density, on the other hand, is a measurement of where the electrons actually are (or aren’t) on a species, and those charges can be fractional or partial charges.

First of all, the charge itself is very real. The ions NH4+ , HO-, H3O+ and so on actually do bear a single charge. The thing to remember is that from a charge density perspective, that charge might be distributed over multiple atoms.
Take an ion like H3O+, for example. H3O *does* bear a charge of +1,

However, if one thinks about where the electrons are in H3O+, one realizes that oxygen is more electronegative than hydrogen, and is actually “taking’ electrons from each hydrogen. If you look at an electron density map of H3O+ , one will see that the positive charge is distributed on the three hydrogens, and the oxygen actually bears a slight negative charge. There’s a nice map here.

When we calculate formal charge for H3O+, we assign a charge of +1 to oxygen. This is for book keeping reasons. As a book-keeping device, it would be a royal pain to deal with fractions of charges like this. So that’s why we calculate formal charge and use it.

Sometimes it does accurately depict electron density. For example, in the hydroxide ion, HO- , the negative charge is almost all on the oxygen.

If you have a firm grasp of electronegativity then it becomes less confusing.

This really helped for neutral covalent molecules. However, I’m having trouble applying this technique for molecules with an overall charge other than 0. For instance, in (ClO2)- , the formal charge of Cl should be 1. However, with your equation the charge should be 0. With the conventional equation, the charge is indeed 1.

I’d appreciate it if you replied sooner rather than later, as I do have a chemistry midterm on Friday. I’m quite confused with formal charges :)